Chem 111 exam answers 2012

Test A5 Name _________________________ Chemistry 11100 Section ____________ FINAL EXAM TA ___________________________ Wednesday, 8:00 AM Total Points = 300 December 12, 2012 Directions: 1. Each student is responsible for following directions. Read this page carefully. 2. Write your name and other requested information on this page and on the separate answer sheet. 3. CODE your name on the answer sheet using an ordinary (#2) pencil. 4. CODE your correct 10-digit identification number (PUID) on the answer sheet. THIS IS VERY IMPORTANT! 5. CODE your section number on the answer sheet. Please use all four digits, 0034, 0035, 0036, etc. This is also very important! 6. CODE the test number shown in the upper right-hand corner on the answer sheet in the block labeled “Test/Quiz Number”. This is Test A5. 7. Put all calculations on the examination pages. DO NOT PUT ANY EXTRA MARKS ON THE COMPUTER ANSWER SHEET! 8. This exam consists of 43 multiple-choice questions worth 6.97 points each. Choose the one best or correct answer for each question and write it both on your exam paper and on the computer answer sheet. The computer answer sheet is the only one that will be graded! 9. This exam consists of 13 pages plus a page of Useful Information, Solubility Rules, a Periodic Table and a sheet of scratch paper. Please check to be sure that you have them all! END OF EXAM 1) Please make sure that you have entered 43 answers on your scan sheet. 2) Make sure that you have entered your name, ID number, and lab section number (4 digits). 3) You MUST turn the scan sheet in to your TA before leaving the exam! KEEP YOUR ANSWERS AND WORK COVERED TO PROTECT THE INTEGRITY OF YOUR WORK!! CHM 11100 – Final Exam --Test #A5 – 12/12/12 _______ 1. Convert 0.458 kg to mg. –3 (a) 4.58 × 10 mg 1 (b) 4.58 × 10 mg 4 (c) 4.58 × 10 mg 5 (d) 4.58 × 10 mg 6 (e) 4.58 × 10 mg _______ 2. Naturally occurring copper consists of copper-63 (62.9296 amu), and copper-65 (64.9278 amu). Using the relative atomic mass from the periodic table, which of the following is the best estimate of the percent abundance of the two isotopes of copper? (a) (b) (c) (d) (e) _______ 3. Which set of ions is formed when (NH4)2CO3 dissolves in water? (a) (b) (c) (d) (e) _______ 10% copper-63 and 90% copper-65 25% copper-63 and 75% copper-65 50% copper-63 and 50% copper-65 75% copper-63 and 25% copper-65 90% copper-63 and 10% copper-65 2 NH4+ , 1 CO32– 1 (NH4)2 , 1 CO3 1 (NH4)2+ , 1 CO3í 2 Ní, 8 H+, and 1 CO32– 2 NH4+, CO3– 4. What is the formula for phosphoric acid? (a) H2PO4 (b) H2PO3 (c) H3PO4 (d) H3PO3 (e) H4PO4 _______ 5. In lab you measure the mass of a metal sample and find it to be 39.22 g. You place it in a graduated cylinder and the water level rises from 15.50 mL to 21.00 mL. What is the density of the solid? (a) (b) (c) (d) 1.87 g/mL 2.53 g/mL 6.53 g/mL 7.13 g/mL 2 CHM 11100 – Final Exam --Test #A5 – 12/12/12 _______ 6. Which of the following best describes HCl when dissolved in water? (a) (b) (c) (d) _______ 7. Which set of formulas is correct for the compounds ammonia, nitric acid, and nitrous acid, respectively? (a) (b) (c) (d) (e) _______ strong electrolyte weak electrolyte nonelectrolyte HCl does not dissolve in water NH3, HN, HNO NH4, HNO2, HNO3 NH3, HNO2, HNO3 NH3, HNO3, HNO2 NH4, HNO3, HNO2 8. How many sodium ions are in 125 g of Na2SO4? 24 (a) 1.06 × 10 Na+ ions 23 (b) 5.30 × 10 Na+ ions 24 (c) 1.37 × 10 Na+ ions (d) 1.76 Na+ ions –24 (e) 1.46 × 10 _______ Na+ ions 9. Aluminum reacts with nitric acid according to the following balanced chemical equation. How many moles of hydrogen gas can be produced from 8 moles of aluminum? (Al = 26.98 g/mole; HNO3 = 63.09 g/mole; H2 = 2.016 g/mol) 2 Al(s) + 6 HNO3 (aq) ˧ 2 Al(NO3)3 (aq) + 3 H2 (g) (a) (b) (c) (d) 8 12 18 24 _______ 10. _____ ion is an example of a polyatomic anion. (a) (b) (c) (d) (e) calcium oxide carbonate ammonium chloride 3 CHM 11100 – Final Exam --Test #A5 – 12/12/12 _______ 11. Molarity is defined as (a) (b) (c) (d) (e) grams of solute per liter of solution. moles of solvent per liter of solvent. moles of solvent per liter of solute. moles of solute per liter of solvent. moles of solute per liter of solution. _______ 12. A 200.0 g sample of river water contains 6.50 mg of lead. How many parts per million (ppm) of lead are in the sample? (a) (b) (c) (d) -2 3.25 × 10 ppm 6.50 ppm 21 ppm 32.5 ppm _______ 13. In the chemical equation shown below, what do the coefficients mean? H2 + I2 ĺHI (a) 1 mole of hydrogen reacts with 1 mole of iodine to give 2 moles of hydrogen iodide. (b) 1 atom of hydrogen reacts with 1 atom of iodine to give 2 molecules of hydrogen iodide. (c) 1 molecule of hydrogen reacts with 1 molecule of iodine to give 2 molecules of hydrogen iodide. (d) All of the above answers, (a), (b), and (c), are correct. (e) Only (a) and (c) are correct. _______ 14. Nitrogen triiodide decomposes to give nitrogen and iodine according to the following balanced equation. 2NI3 (s) ˧ N2 (g) + 3I2 (s) How many grams of reactant, NI3, would be required to produce 2.53 g iodine? (NI3 = 394.71 g/mol; N2 = 28.02 g/mole; I2 = 253.8 g/mole) (a) (b) (c) (d) 1.63 g 2.62 g 3.93 g 5.90 g 4 CHM 11100 – Final Exam --Test #A5 – 12/12/12 _______ 15. The correct answer for the addition 102.5 mL + 6.57 mL is: (considering significant figures) (a) (b) (c) (d) 109.07 mL 109 mL 109.0 mL 109.1 mL _______ 16. If the temperature of water in a freezer decreases from 22°C to –25°C, what is the decrease in temperature in units of degrees Celsius and Kelvin? (a) (b) (c) (d) (e) 47°C, 320 K 47°C, 273 K 47°C, 47 K 3°C, 276 K 3°C, 3 K _______ 17. What is the parent structure (electron geometry) of phosphine, PH3? (a) (b) (c) (d) Tetrahedral Trigonal pyramidal Bent Trigonal planar _______ 18. What is the abbreviated electron configuration for the bromide ion? (a) [Ar] (b) [Kr] 2 6 2 5 2 10 (c) [Ar] 3s 3p (d) [Ar] 4s 4p 5 (e) [Ar] 4s 3d 4p 5 CHM 11100 – Final Exam --Test #A5 – 12/12/12 _______ 19. Calcium carbonate, CaCO3, is often used in commercial antacids. It acts to reduce the acidity in the stomach by neutralizing stomach acid, which is mostly HCl, by the following reaction: CaCO3 (s) + 2 HCl (aq) ˧ CaCl2 (aq) + CO2 (g) + H2O (l) What mass of CaCO3 is needed to neutralize 0.0550 mol HCl? (CaCO3 = 100.09 g/mole; HCl = 36.461 g/mol) (a) (b) (c) (d) –2 17.55 x 10 g 2.75 g 5.50 g 11.0 g _______ 20. When comparing a 10.00 g sample of iron (Fe) with a 10.00 g sample of lead (Pb) what is true? (a) (b) (c) (d) Each sample has the same number of atoms. There are more iron atoms than lead atoms. There are more lead atoms than iron atoms. The lead is heavier than the iron, and therefore there would be more atoms. _______ 21. Which set of elements below contains, respectively, an alkaline earth metal, a noble gas, and a metalloid? (a) (b) (c) (d) (e) Na, Ar, Si Ba, O, As Ti, Cl, Pb Bi, Kr, B Mg, Ne, Ge _______ 22. Which combination of formula and name is incorrect? (a) (b) (c) (d) (e) K+ = potassium ion I- = iodide ion Mg+ = magnesium ion S2- = sulfide ion N3- = nitride ion 6 CHM 11100 – Final Exam --Test #A5 – 12/12/12 _______ 23. For each of the following pairs of ions, indicate which pair would form a precipitate when solutions containing these ions are mixed. 2+ – (a) Fe and I + – + 2– (b) Ag and NO3 (c) Na and SO4 + – (d) K and OH 2+ 2– (e) Ca and CO3 _______ 24. Iron and oxygen react to form iron(III) oxide. 4 Fe (s) + 3 O2 (g) ˧ 2 Fe2O3 (s) What is the limiting reactant in a mixture of 6.0 moles of Fe and 6.0 moles of O2? (a) O2 (b) Fe (c) Fe2O3 (d) Impossible to determine _______ 25. 20.0 g NH3 and 20.0 g CO2 were reacted according to the following balanced chemical equation. 2 NH3 + CO2 o CH4N2O + H2O How many grams of urea (CH4N2O) are produced? (NH3 = 17.034 g/mole; CO2 = 44.01 g/mole; CH4N2O = 60.062 g/mole) (a) (b) (c) (d) 27.3 g 35.3 g 62.6 g 70.5 g 7 CHM 11100 – Final Exam --Test #A5 – 12/12/12 _______ 26. What is the molecular shape of the species below that has 3 bonded atoms and 1 unshared pair of electrons? (a) (b) (c) (d) (e) octahedral trigonal bipyramidal trigonal pyramidal tetrahedral trigonal planar _______ 27. Which of the following molecules is polar? _______ 28. The frequency of microwave radiation in a common microwave oven that you might have in 9 –1 your kitchen is 2.46 x 10 s . What is the wavelength of this radiation? 17 (a) 1.20 x 10 m 17 (b) 7.38 x 10 m –1 (c) 1.22 x 10 m 8 (d) 3.00 x 10 m 9 (e) 2.46 x 10 m 8 CHM 11100 – Final Exam --Test #A5 – 12/12/12 _______ 29. Dissolving calcium chloride in water is an exothermic process. Which of the following statements is correct about the temperature of the resulting calcium chloride solution? (a) The temperature of the resulting solution will be lower than the initial temperature of the pure water. (b) The temperature of the resulting solution will be higher than the initial temperature of the pure water. (c) The temperature of the resulting solution will be the same as the initial temperature of the pure water. (d) The temperature of the water is not dependent on what is dissolved in it. (e) There is not enough information about this process to determine how the temperature will change. _______ 30. Mixing a solution of lithium sulfate with a solution of barium chloride results in the formation of solid, white barium sulfate, as shown in the equation below. What is the net ionic equation for this reaction? Li2SO4(aq) + BaCl2 (aq) o 2 LiCl (aq) + BaSO4(s) (a) Li2SO4(aq) + BaCl2 (aq) o 2 LiCl (aq) + BaSO4(s) + 2– 2+  + – + 2– 2+  + – 2+ 2– 2+ 2– (b) 2 Li (aq) + SO4 (aq) + Ba (aq) +2 Cl (aq) o 2 Li (aq) + 2 Cl (aq) + Ba (s) +2 SO4 (s) (c) 2 Li (aq) + SO4 (aq) + Ba (aq) +2 Cl (aq) o 2 Li (aq) + 2 Cl (aq) + BaSO4(s) (d) Ba (aq) + SO4 (aq) o BaSO4(s) +  (e) Li (aq) + Cl (aq) o LiCl(aq) _______ 31. A compound has a melting point of 545 oC and dissolves well in water. What true about about the bonding in the compound and the reasoning to support this conclusion? (a) (b) (c) (d) It is ionic because it has a high melting point and dissolves in water. It is covalent because it has a high melting point and dissolves in water. It is ionic because it has a low melting point and is insoluble in water. It is covalent because it has a low melting point and is insoluble in water. _______ 32. Identify the most polar bond. (a) (b) (c) (d) O-Cl C-Cl H-Cl Cl-Cl 9 CHM 11100 – Final Exam --Test #A5 – 12/12/12 _______ 33. Identify the correct Lewis structure for CH2Cl2. _______ 34. Copper metal dissolves in dilute nitric acid HNO3 according to the following equation: 3Cu(s) + 8HNO3(aq) o 3Cu(NO3)2(aq) + 2NO(g) + 4H2O(aq) How many grams of NO(g) can be produced from 10.0 grams of Cu? (a) (b) (c) (d) 7.08 g 4.72 g 3.15 g 0.105 g 10 CHM 11100 – Final Exam --Test #A5 – 12/12/12 _______ 35. Indicate the parent structure (electron geometry) and the molecular shape of the species below. (a) (b) (c) (d) (e) Parent Structure or Electron Geometry Tetrahedral Trigonal planar Trigonal planar Linear Trigonal planar Molecular Shape trigonal pyramidal trigonal planar linear linear bent _______ 36. Which of the following molecules or ions exhibits resonance? (a) N2 (b) O3 (c) CO2 (d) SO3 2– ______ 37. A 25.0 mL sample of an HCl solution is placed in a flask with a few drops of phenolphthalein indicator. This solution is titrated with NaOH. If 32.6 mL of a 0.185 M NaOH solution is needed to reach the end point, what is the concentration of the HCl solution? –4 (a) 2.41 × 10 M –3 (b) 6.03 x 10 M (c) 0.241 M (d) 0.142 M _______ 38. Which molecule, CF4 or CCl2F2, is the most soluble in water and why? (a) CF4 because its molar mass is lighter. (b) CF4 because it is tetrahedral shaped. (c) CCl2F2 because its molar mass is heavier. (d) CCl2F2 because it is polar. 11 CHM 11100 – Final Exam --Test #A5 – 12/12/12 _______ 39. Among these four ions Mg2+, Na+, F- O2- which has the smallest ionic radius and why? 2– (a) O because it has the most electrons. – (b) F because it is isoelectronic with Ne. + (c) Na because it has the smallest atomic radius. 2+ (d) Mg because it has the greatest number of protons. _______ 40. Which color of visible light has the most energetic photons? (a) (b) (c) (d) (e) Red, 650 nm Yellow, 570 nm Green, 510 nm Blue, 475 nm Violet, 400 nm _______ 41. What did you determine (or measure) about the wine you analyzed in the “Analysis of Wine” lab? I. II. III. IV. V. (a) (b) (c) (d) pH Alcohol content Acid content (or concentration) Density Conductivity I, II, III I, II, III, IV I, II, III, IV, V I, II, III, V _______ 42. In both the “Analysis of Wine” lab and “Training Lab 2” you performed titrations. In both experiments you added an indicator (phenolphthalein) to the solution in an Erlenmeyer flask which contained an acid. Why did you add the indicator? (a) (b) (c) (d) To buffer the solution. To identify when all the acid had been neutralized by the base. To hold the pH of the acid constant. To measure the initial volume of the acid. 12 CHM 11100 – Final Exam --Test #A5 – 12/12/12 Absorbance  at  640  nm   _______ 43. In a lab to determine the concentration of red dye #40 in Purplesaurus Rex Kool-Aid a student made the following calibration curve. The trend line is on the plot. Absorbance  versus  Concentration   0.8 0.7 0.6 0.5 0.4 0.3 0.2 0.1 0 0.00E+00 y  =  1210.2x  -­‐  0.0187   R²  =  0.9874   1.00E-­‐04 2.00E-­‐04 3.00E-­‐04 4.00E-­‐04 5.00E-­‐04 6.00E-­‐04 7.00E-­‐04 Concentration  of  Red  Dye  #40  (M)   The absorbance of the Kool-Aid solution was 0.41. What was the concentration of red dye #40 in the solution? –4 (a) 3.5 x 10 M –4 (b) 3.2 x 10 M (c) 0.41 M (d) 496 M 13 CHM 11100 – Final Exam --Test #A5 – 12/12/12 Useful Information   % Recovery %  Yield =   TK mass of material recovered u 100%   mass of material started with Actual   ×  100 Theoretical ToC  273.15 ToF 1.8(ToC )  32 MiVi = MfVf 6 1 ppm = 1 g/ 1 x 10 g = 1 mg/1 L 23 Avogadro’s number: 1 mole = 6.022 x 10 formula units 4.184 J = 1 cal 1000 cal = 1 Cal ‫ ݉ = ݍ‬  ×  ‫ ܥ‬  ×   οܶ ܿ =  ߣߥ                                          ‫ = ߥ݄ = ܧ‬  c = 3 x 108 m/s ݄ܿ ߣ h = 6.626 x 10-34 J s 14 CHM 11100 – Final Exam --Test #A5 – 12/12/12 Solubility rules from Table 5.4: Rules used to predict the solubility of ionic compounds. Ions + Rule + + Na , K , NH4 (and all other alkali metal ions) Most compounds of alkali metal and ammonium ions are soluble. NO3-, CH3CO2- All nitrates and acetates are soluble. SO42- Most sulfates are soluble. Exceptions are BaSO4, SrSO4, PbSO4, CaSO4, Hg2SO4, and Ag2SO4. Cl-, Br-. I- Most chlorides, bromides, and iodides are soluble. Exceptions are AgX, Hg2X2, PbX2, and HgI2 (X = Cl, Br, or I). Ag+ Silver compounds except AgNO3 and AgClO4 are insoluble. AgCH3CO2 is slightly soluble. O2-, OH- Oxides and hydroxides are insoluble. Exceptions are alkali metal hydroxides, Ba(OH)2, Sr(OH)2, and Ca(OH)2 (somewhat soluble) S2- Sulfides are insoluble. Exceptions are compounds of Na+, K+, NH4+, Mg2+, Ca2+, Al3+, and Ni2+ CO32-, PO43-, SO32- Most carbonates, phosphates, and sulfites are insoluble. Exceptions are compounds of Na+, K+ , and NH4+ 15 CHM 11100 – Final Exam --Test #A5 – 12/12/12 Key for A5 and E1 1.) 2.) 3.) 4.) 5.) 6.) 7.) 8.) 9.) 10.) 11.) 12.) 13.) 14.) 15.) 16.) 17.) 18.) 19.) 20.) 21.) 22.) 23.) 24.) 25.) 26.) 27.) 28.) 29.) 30.) 31.) 32.) 33.) 34.) 35.) 36.) 37.) 38.) 39.) 40.) 41.) 42.) 43.) D D A C D A D A B C E D E B D C A B B B E C E B A C B C B D A C D C E B or D C D D E A B A 16